Enfield Public Schools

Enfield Public Schools

Advanced (AP/UCONN) Chemistry

(0297)

Curriculum Writers:

Patrick Smith

William Schultz

November 2007

Lab Safety
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. Basic safety rules must be followed in the Advanced Chemistry laboratory. / ·  Follow teacher’s instructions and notify teacher when there is a problem.
·  Use safety equipment properly.
·  Handle chemicals and equipment safely.
·  Dispose of chemical waste and broken glass properly.
·  Keep lab station clean.
·  Notify teacher if accident occurs.
·  Read hazard labels or Material Safety Data Sheets (MSDS) for chemicals used.
Matter & Measurements
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. All forms of matter can be classified.
2. The proper handling of measurements is necessary in order to maintain the accuracy and precision of an experiment.
3. All substances have unique chemical and physical properties that can be used for identification. / ·  Classify a sample of matter as either an element, compound, homogenous mixture, or heterogeneous mixture.
·  Make calculations with measurements using and reporting the correct number of significant figures.
·  Distinguish between extensive and intensive properties.
Atoms, Molecules, and Ions
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. The atomic theory underlies all of chemistry. Various experiments including the Cathode Ray Tube experiment and the Gold Foil Experiment led to our modern view of the atom.
2. The atom is composed of protons, neutrons, and electrons.
3. The Periodic Table contains different regions of elements.
4. Besides existing in their elemental state, atoms can form molecules, ions, and ionic compounds.
5. Chemical formulas are used to express the chemical composition of compounds.
6. There is a systematic way of naming compounds based on their composition. / ·  Explain how various experiments led to our modern view of the atom.
·  Identify and list the properties of the three subatomic particles.
·  Classify an element according to the regions on the periodic table. (e.g., alkali metals, alkali earth metals, transition metals, metalloids, halogens, noble gases)
·  Write chemical formulas for molecules and ionic compounds.
·  Distinguish between the different types of formulas. (e.g., molecular, structural, condensed structural, empirical.)
·  Name ionic and binary molecular compounds based on their chemical formulas.
·  Name acids based on their chemical formula.
Mass Relations in Chemistry; Stoichiometry
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. The atomic mass unit is defined as one-twelfth the mass of a Carbon-12 isotope.
2. The mole is defined as 6.022 X 1023 particles of a substance and is a critical tool in quantitative chemistry.
3. Using analytical data, the empirical and molecular formula of a compound can be calculated.
4. Based on the Law of Conservation of Mass, chemical reactions contain various mass relationships that can be found and calculated. / ·  Use the mole concept and Avogadro’s number to convert between mass, moles, atoms, and molecules.
·  Determine the empirical formula of a compound given the mass per cent of the elements or analytical data.
·  Determine the molecular formula of a compound given its empirical formula and molar mass.
·  Balance chemical equations and make stoichiometric calculations using those equations.
Reactions in Aqueous Solutions
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. Molarity (moles solute / Liters of solution) is a common unit used to express the concentration of an aqueous solution.
2. Three critically important reactions in aqueous solution are precipitation reactions, acid-base reactions, and oxidation-reduction reactions. / ·  Do various concentration calculations using molarity.
·  Write net ionic equations for precipitation reactions.
·  Perform stoichiometric calculations for precipitation reactions.
·  Distinguish between strong and weak acids and bases.
·  Write acid-base reactions.
·  Perform titration calculations for precipitation, acid-base, and oxidation-reduction reactions.
·  Determine the oxidizing and reducing agents in an oxidation-reduction reaction using oxidation numbers.
·  Balance oxidation-reduction equations.
Gases
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. The Ideal Gas Law relates the pressure, volume, number of moles, and temperature of a sample of gas.
2. The Ideal Gas Law can be used to make stoichiometric calculations with gases.
3. Calculations involving gas mixtures can be performed using Dalton’s Law of Partial Pressures and the concept of the mole fraction.
4. The Kinetic Theory of Gases explains the behavior of a gas in terms of small particles that are in constant, random motion.
5. Under certain conditions, real gases deviate from the behavior predicted by the Ideal Gas Law. / ·  Perform Ideal Gas Law calculations. (e.g., calculate the volume of a gas given its temperature, pressure, and number of moles.)
·  Use the Ideal Gas Law to make calculations involving chemical equations. (e.g., calculate the volume of product formed when a given volume of reactant is given.)
·  Apply Dalton’s Law of Partial Pressures and/or the concept of the mole fraction of a gas in order to calculate the partial pressures of component gases in a mixture.
·  List the assumptions to the Kinetic Theory of Gases and explain how it relates to the Ideal Gas Law.
·  Identify the conditions under which a gas deviates from Ideal behavior.
Electronic Structure and the Periodic Table
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. The wavelengths of light given off by an excited atom are unique to each element and are a result of the energy states of the electrons in that element’s atoms.
2. Niels Bohr developed a model which explains the atomic spectrum of the hydrogen atom using the concept of quantum energy levels.
3. The specific energy state of an electron can be described by its four quantum numbers.
4. Electron configurations are used to describe the arrangement of electrons in an atom.
5. The Atomic Radius, Ionic Radius, Ionization Energy, and Electronegativity of the elements change in periodic fashion. / ·  Explain the idea of atomic spectra and how they can be used to identify an element.
·  Use the equation c = λν to make calculations involving the wavelength and frequency of a wave of light.
·  Use the equation Ep = hν to calculate the energy of a photon given a certain frequency of light.
·  Use the Bohr model of the atom to calculate the wavelength of light emitted or absorbed from certain energy level transitions in the hydrogen atom.
·  Explain the concept of quantum energy levels in the atom.
·  Explain the Quantum Mechanical model of the atom and how it differs from Bohr’s model of the hydrogen atom.
·  Use quantum numbers to (m, l, ml, ms ) to describe the energy state of an electron.
·  Write electron configurations and orbital diagrams for various elements and their ions.
·  Compare various properties of elements based on their position on the periodic table. (e.g., atomic radius, ionic radius, ionization energy, electronegativity.)
Covalent Bonding and Shapes
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. Lewis structures are used to illustrate the electron distribution in covalently bonded molecular compounds. The Octet Rule is the underlying principle for Lewis Structures.
2. The VSEPR Theory and Lewis structures can be used to predict the three dimensional geometry of molecules.
3. Polarity is the uneven distribution of electrons in molecules due to geometric factors and electronegativity.
4. Hybridization of atomic orbitals is used to explain the geometry of molecules. / ·  Write Lewis Dot diagrams of various molecules.
·  Predict the geometry and polarity of a molecule based on its Lewis Dot diagram.
·  Explain the concept of Resonance with regards to electron distribution.
·  Determine whether molecules are polar.
·  State the hybridization of the orbitals in a given molecule and relate to double and triple bonds.
Thermochemistry
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. In all processes, heat flows from one substance to another. The extent of heat flow can be calculated per q = mcΔT.
2. Heat flow can be measured using the concept of calorimetry.
3. Enthalpy (H) is the heat that flows in or out of a chemical substance as a result of a reaction that that substance undergoes.
4. Balanced chemical equations can be written which include the enthalpy change (ΔH) for the reaction. These are called thermochemical equations.
5. The enthalpy changes associated with the formation of elements and compounds are tabulated and can be used to calculate the enthalpy change for other reactions. / ·  Perform heat flow calculations.
·  Distinguish between positive and negative heat flow in relation to the system and surroundings.
·  Distinguish between exothermic and endothermic reactions.
·  Explain how a calorimeter works to obtain thermochemical data.
·  Explain the concept of Enthalpy and how it relates to heat flow at constant pressure.
·  Write and make calculations using thermochemical equations.
·  Use the Enthalpies of Formation of reactants and products in order to calculate the enthalpy change for a reaction.
Liquids and Solids
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. Whenever a liquid is present, a Liquid-Vapor Equilibrium exists. The pressure of the vapor at this equilibrium increases rapidly as temperature increases toward the liquid’s boiling point.
2. Phase Diagrams illustrate which physical state (solid, liquid, or gas) is present at any given pressure and temperature combination.
3. In Molecular Substances, the Intermolecular Forces (i.e. forces between molecules) are responsible for the properties of the substances.
4. In Network Covalent, Ionic, & Metallic Solids, intramolecular forces (i.e. forces within molecules, so chemical bonds) are responsible for the properties of substances.
5. Crystal substances exist with specific structures that relate to their formulas and properties. / ·  Explain the concept of vapor pressure and its relationship to temperature and boiling point.
·  Use the Clausius-Clapeyron equation to perform various vapor pressure calculations.
·  Read and interpret phase diagrams.
·  Identify the intermolecular forces present in a given substance.
·  Compare the boiling points of substances based on their intermolecular forces.
·  Determine what type of bonding is present in compounds.
·  Distinguish between the various types of unit cells that make up crystal solids. (e.g., simple cubic, body-centered, face-centered.)
·  Determine the number of atoms in each type of unit cell.
·  Perform ionic radius calculations using the concept of the unit cell.
Solutions and Colligative Properties
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. The amount of solute dissolved in a solvent can be measured using various concentration units (molarity (M), molality (m), mol fraction (X), and mass %).
2. Various principles of solubility are utilized to predict whether solutes will dissolve in solvents.
3. Certain properties of solutions depend only on the amount of solute dissolved, not on the type of solute. These are called Colligative Properties. Freezing point depression, boiling point elevation, vapor pressure lowering, and raising of the osmotic pressure are 4 colligative properties. / ·  Convert between each of the concentration units.
·  Predict solubilities of various solutes in various solvents.
·  Perform calculations based on colligative properties.
Rates of Reactions
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. Specific techniques are used to
measure and express reaction rates.
2. Reaction rate & concentration are related by the Rate Law.
3. Different reactants have different affects on reaction rates as summarized by the order of reaction.
4. Identify how reactant concentration varies with time for 0, 1st, and 2nd order reactions.
5. Molecular principles underlie our mathematical models for reaction rate.
6. Reaction rates vary with temperature.
7. Catalysts (both homogeneous and heterogeneous) increase reaction rates
8. Reaction mechanisms explain the underlying molecular principles. / ·  Be able to mathematically relate the reaction rate to the rate of change in the concentrations of the reactants and products.
·  Give an example of increased concentration causing an increase in reaction rate (compression of fuel-air mixtures in engines, 3% versus 30% hydrogen peroxide, metal in dilute versus concentrated acid).
·  Determine orders of reactions from data on rates versus concentrations
·  Calculate reactant concentration at any given time after a reaction begins.
·  Explain the collision and activation energy principles.
·  Use the Arrhenius Equation to calculate rate constants at different temperatres.
·  Give an example of a catalyst (iodide with hydrogen peroxide, enzymes, catalytic converters.
·  Evaluate a proposed reaction mechanism and determine if it is appropriate based on comparison to the rate law for the reaction.
Gaseous Equilibrium
Learner outcomes / Sample indicators/assessments
The student(s) will know (how to): / The student(s) will be able to:
1. The N2O4-NO2 system is a simple introduction to the concept of equilibrium.
2. The equilibrium constant expression is a mathematical expression that relates reactant and product concentrations for a system at equilibrium.
3. The equilibrium constant (K) can be determined by analysis of concentrations in reaction mixtures.
4. The equilibrium constant (K) can be used to predict the direction and the extent of reactions.
5. The equilibrium constant (K) can be used to determine the extent of reactions.
6. Changes in conditions imposed on an equilibrium system will result in a shift in the equilibrium position (Le Chatelier’s Principle). / ·  Describe the shifts in quantities of N2O4 and NO2 present in the equilibrium system as the temperature or pressure is changed.
·  Write the equilibrium constant expression for any chemical reaction.